The Periodic Table
Objectives
1. Describe the contributions of Dmitri Mendeleev and Henry Moseley to the
development of the modern periodic table.
2. Describe the arrangement of elements on the periodic table.
3. Define "group" and "period" and describe how the properties of elements change in
groups and periods.
4. Locate the metals, metalloids and nonmetals on the periodic table.
5. Describe the metals, metalloids and nonmetals.
6. State the periodic law.
7. State which subatomic particle plays the greatest role in determining the chemical and
physical properties of the different elements.
8. Describe the noble gases, representative elements, transition metals and inner transition
metals in terms of their electron configurations.
9. Identify the s, p, d, and f blocks on the periodic table.
10. Describe and explain the groups and periodic trends in atomic size, ionization energy,
electron affinity, ionic size and electronegativity.
11. Define "ionization energy," " electron affinity," and "electronegativity."
12. Name and describe three forces affecting the periodicity of elements.
13. Describe the following groups and state several uses of each : noble gases, alkali
metals, alkaline earth metals, aluminum group, carbon group, nitrogen group, oxygen
group and the halogens.
14. Name and explain the accepted theory about the origin of the elements which make
up the universe.
1. Dmitri Mendeleev (1834-1907) was a Russian chemist who was one of the first persons who
tried to put the known elements (about 70 during Mendeleev's time) in a chart which showed a
pattern of properties. Mendeleev arranged the known elements in vertical columns by increasing
atomic mass and noticed patterns in their properties. From his arrangement Mendeleev could
predict the properties of elements which had not yet been discovered.
Henry Moseley (1887-1915) arranged the elements in order of increasing atomic number
(positive charge). This is the way that the modern periodic table is arranged today.
2 and 3. The elements are arranged in order of increasing atomic number across the periods
(horizontal rows). The periods are arranged so that the elements in the vertical columns (groups
or families) have similar properties. This causes the properties of elements to change as you
move horizontally from group to group across a period.
Note that their are 7 periods - one for each energy level (principle quantum number).
4 and 5. See the periodic table in your reference Tables
6. The periodic law states that when the elements are arranged according to increasing atomic
number there is a periodic pattern in their physical and chemical properties.
7 and 8. Although the periodic table is arranged by increasing atomic number (which indicates
the number of protons), the electrons configuration is what really determines the physical and
chemical properties of the elements. The periodic table can be divided into four groups based on
electron configuration :
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The Noble gases (Group 0) - have their outermost s and p orbitals filled which
creates a stable and non-reactive (inert) element.
The representative elements - Group A elements - have their s and p orbitals being
filled. These include :
o Group 1 - Li, Na, K etc. - all very reactive with one electron in the outer s
orbital
o Group 2 - Be, Mg, Ca etc. - all quite reactive with 2 electrons filling their
outer s orbital
o Group 13 - Aluminum group - 3 electrons in outer energy level (2s and 1p)
properties vary from metallic to metalloid
o Group 14 - Carbon group - 4 electrons in outer energy level (2s and 2p) properties vary from nonmetallic to metalloid to metallic down the group
o Group 15 - Nitrogen group - 5 electrons in outer energy level (2s and 3p) properties vary from nonmetallic to metalloid to metallic
o Group 16 - Oxygen group - 6 electrons in outer energy level (2s and 4p) properties vary from nonmetallic to metalloid
o Group 17 - Halogens - all have 7 electrons in the outer energy level (2s
and 5p) - properties vary from nonmetallic to metalloid. Very reactive due
to the outer energy level being almost filled.
The transition metals - elements whose d orbitals are being filled - found in the
"d-block." These are also called the Group B elements
The Inner transition metals - These are the Lanthanide and Actinide series,
element whose f orbitals are being filled.
9. The s, p, d, and f groups can be identified on the diagram below. The f block (inner transition
metals) is usually shown separated and below the rest of the table.
10 and 11. Periodicity is the property of having periodic properties. The periodic table shows
periodicity in the following properties :
M E R I
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Atomic size - atoms of elements tend to increase as you go down a group (due to
a greater number of energy levels) and atoms of elements tend to decrease in size
across a period (greater positive nuclear charge which draws in electrons -energy
levels are constant across a period).
Ionization energy - the energy required to remove an electron from the gaseous
state of the atom. Ionization energy decreases as you go down a group due to the
outer electrons being further from the positive charge of the nucleus and being
shielded from the nucleus' positive charge by the inner energy levels. Ionization
energy increases as you move across a period. This is due to the increase in
nuclear charge without the increase in number of energy levels.
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Electron affinity - this is the energy change associated with the addition of an
electron to a gaseous atom. This trend is not as consistent as the others, but in
general electron affinity decreases down a group and increases across a period.
Ionic size - The size of ions increases as you go down a group, and decreases as
you move across a period for the metals and for the nonmetals, for the same
reasons as atomic size. However, the metallic ions are positive (have lost
electrons, which make up the space or size of the atom) and are much smaller than
the negative nonmetallic ions (have gained electrons which create the volume of
atoms or ions).
Electronegativity -the tendency of an atom to gain an electron(s) when combining
with another element. Electronegativity decreases down a group (due to shielding)
and increases as you move across a period (due to the increase in nuclear charge).
12. The three major forces affecting periodicity are :
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Nuclear charge - the greater the number of protons in the nucleus, the greater the
positive charge and the stronger the electrons are held.
Shielding - the effect of inner energy levels reducing the strength of the nuclear
charge on the electrons in the outer energy levels.
Electron configuration - atoms are most stable when their outer orbitals are filled
(especially the s and p orbitals). This causes the Noble gases to be inert.
13. The representative groups of elements :
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Alkali metals - Group 1, lithium, sodium, potassium, rubidium, cesium and
francium
 very reactive (one electron away from a filled s and p orbital)
 low density
 low melting point
 good electrical conductivity
 react with water to form strong bases (sodium hydroxide, lithium
hydroxide etc.)
Alkaline earth elements - Group 2, beryllium, magnesium, calcium, strontium,
barium and radium
 very reactive (2 electrons away from a filled s and p orbital)
 react with water to form hydroxides
 used to form metal alloys
Aluminum group - Group 13 - 3 electrons in outer energy level (2s and 1p)
properties vary from metallic to metalloid
 aluminum is the most useful metal of this group being lightweight and
strong to make boats, aircraft etc.
Group 14 - Carbon group - 4 electrons in outer energy level (2s and 2p) properties vary from nonmetallic to metalloid to metallic down the group
 diamond and graphite are forms of pure carbon
 silicon and germanium are semiconductors used in electronics
 tin and lead are useful metals
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Group 15 - Nitrogen group - 5 electrons in outer energy level (2s and 3p) properties vary from nonmetallic to metalloid to metallic
 nitrogen and phosphorus are elements necessary to form proteins and
nucleic acids in living things
Group 16 - Oxygen group - 6 electrons in outer energy level (2s and 4p) properties vary from nonmetallic to metalloid
 oxygen is the most abundant element on the earth
 sulfur has many industrial uses (sulfuric acid is the most widely used
industrial chemical)
Group 17 - Halogens - all have 7 electrons in the outer energy level (2s and 5p) properties vary from nonmetallic to metalloid. Very reactive due to the outer
energy level being almost filled.
 iodine is used as an antiseptic
 chlorine is a bleaching agent and disinfecting agent
 fluorine, as the fluoride ion, is used to maintain the health of our teeth
 fluorine is used to make Teflon
Group 18 or O: Noble Gases , helium, neon, argon, krypton, xenon and radon
 inert (unreactive) because of stable electron configuration (filled s and p
orbitals)
 helium is used in weather balloons
 helium and neon are used to create artificial, unreactive environments (less
soluble than nitrogen and therefore less likely to cause the bends
 other noble gases are used to create unreactive environments in flashbulbs
or aluminum welding
33. The accepted theory which explains the origin of the elements and all matter in the universe
is the Big Bang theory. It states that the universe began with an explosion of tremendous energy.
This energy was converted into matter, according to Einstein's equation E=mc2. At first all
matter was in the form of quarks. As the universe expanded it cooled allowing matter to
condense and form the lightest elements first and then the heavier elements.